This invention relates to an electrochemical cell used in reducing a metal oxide or a combination of metal oxides to the corresponding metal in which an electric potential is established between an anode and cathode, usually in a molten salt electrolyte. In general, the molten salt electrolyte is selected from one or a combination of the alkali metal halides or the alkaline metal earth halides and the oxide to be reduced is positioned in a container which acts when filled with the metal oxide as a cathode, and the anode is generally some material which is impervious to attack by the products in an anode reaction or by the electrolyte or dissolved materials within the electrolyte.
In general, the reaction is according to the following wherein the base metal is (M) and the oxide to be reduced is (MOx), both contained in a molten salt electrolyte, the electrochemical cell operating according to the overall reaction (1):                               M          ⁢                                           ⁢                      O            x                          →                  M          +                                    x              2                        ⁢                          O              2                                                          [        1        ]            
Although this invention will be explained using an example of spent oxide-based nuclear fuels, which are mostly uranium oxide, the invention is not so limited and pertains to variety of metals such as the actinide and rare earths, except for those metal oxides having sufficiently negative Gibbs free energy of formation such that reduction is not commercially advantageous. Classically this process is accomplished with a two-electrode electrochemical cell having a cathode where the metal oxide is reduced to the metal and a second electrode where oxygen is evolved through the oxidation of oxygen anions, called the anode. The individual electrode reactions are given by [2] for the cathode and [3] for the anode.MOx+2xe−→M+xO−2  [2]20−2→O2+4e−  [3]
Although a number of molten salts, such as CaCl2 can be used for the electrolyte in this process, work has concentrated on lithium chloride. An oxide traditionally has to be present in the electrolyte melt to sustain a significant current density at the oxygen-evolving anode.
Several difficulties with the two-electrode cell have been identified. These include: (1) Cell control for optimum efficiency. It is necessary to maintain the anode potential at the highest possible value to increase cell throughput, but this value must be kept below the potential where excessive chlorine evolution occurs. Likewise, the cathode must be maintained at the lowest practical potential for best throughput, but not so low that metallic lithium is deposited on and in the metal oxide. Maintaining these voltages requires a very stable reference and a good means for constantly adjusting the power supply output. The voltage must be limited to accommodate whichever electrode is slowest at various stages of the reduction process; the current, and hence the rate, is limited correspondingly. (2) Elimination of oxygen-containing compounds from the cell. Any oxygen compounds in the cathode product or the electrolyte that is inevitably carried with the cathode product are deleterious to following operations. These compounds include lithium oxide in the electrolyte and unreduced oxides, such as rare earth oxides, that will be mixed with the cathode product and can not be eliminated without generating excessive chlorine. (3) The demanding performance requirements that this cell places on the anode. The current in a two-electrode cell is highest at the beginning of operation, and decays to a low value at the end. The anode must support the full current of the cell early in the reduction process, and must continue to function without excessive chlorine generation at the end of the process, when it is desirable to reduce the lithium oxide concentration in the electrolyte, preferably to one ppm or less.